Bio 121-999-F/12 LEC-2 - Missouri State...
Transcript of Bio 121-999-F/12 LEC-2 - Missouri State...
8/20/2012
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Interplay of Biology and Chemistry
Here is a link to the video…these beetles are fairly common locally – an amazing adaptation, and a good example of chemistry and physics in biology.
Also look for creationist-evolutionist arguments about these on the internet. Bombardier beetle
Definition of Matter
• Matter has mass and occupies space.
• Mass is a property of matter that causes inertia and weight.
• Matter is composed of several kinds of subatomic particles that combine to make atoms.
Subatomic particles
• Proton has mass of 1 and an electrical charge of +1
• Neutron has mass of 1 and no charge
• Electron has mass near zero and electrical charge of -1
• Opposite charges attract, so protons attract electrons
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Atoms
• Protons and neutrons in small, dense nucleus
• Protons (+) attract electrons (-) which surround the nucleus
• Rutherford model- see p. 22 Brooker
Neutral vs Ionized
• Atoms with equal numbers of protons & electrons are electrically neutral
• Ions are atoms with a net electrical charge
• Cations are positively charged(more protons than electrons)
• Anions are negatively charged(more electrons than protons)
Charge & electrostatic force
• Opposite charges attracteach other and balanceeach other at close range. Like charges repel each other
• When opposite charges are separated, or similar charges are together, they have energy (electrostatic force) static cling or static “fling” ?
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• Molecular shape
• Cell membrane potential & coupled transport processes
• Nerve & muscle action potentials
+/-Charge separation and electrostatic forces lead to…
Elements
• Kinds of atoms, each with unique number of protons (= atomic number)
• Atomic number is indicated by a left subscript. For example: 6C (carbon)
• Periodic table lists the elements and their properties.
Isotopes
• Forms of an element that differ in the number of neutrons in the nucleus
• Atomic mass is indicated by a left superscript, e.g. 14C (carbon-14) or 12C
• Isotopes of an element have different mass but same atomic number and similar chemical properties.
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Electrons
• move around the nucleus in patterns called shells, subshells, and orbitals
• follow rules called “quantum mechanics”
• First shell holds up to 2 electrons. The second and third shells hold up to 8 electrons each…..
Electron orbitals
These patterns have complex shapes (upper row) but are often diagrammed as circles (lower row, below)
Electron configurations of the first 18 elements
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What atoms “want”
1. Full outermost shell (valence shell)
2. No net electrical charge (i.e. equal numbers of protons and electrons)
• “noble gases” have the right atomic numbers do both.
• Other atoms share electrons to fill the valence shell: chemical bonds result
Noble elements- examples
• Helium (2He) has 2 protons, so 2 electrons fill first shell
• Neon (10Ne) has 10 protons, so 10 electrons fill first 2 shells
• Both are chemically unreactive (don’t make bonds with other atoms)
Covalent bonds
• Two or more atoms share electrons in a combined valence (covalent) shell
• Single or double bonds: one or two pairs of electrons may be shared
• Shared electrons bind the atoms together
Note: The blue area represents the shared electrons
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Examples of molecules with covalent bondsnote the 3 different types of diagrams are shown below- all illustrate the same 3 molecules.
Polar covalent bonds
• nonpolar = equal sharing of electrons
• polar = electrons spend more time near one nucleus than the other
• Therefore the charge distribution is “polar” (meaning that there are positive and negative ends)
- +
Note: The blue area represents the shared electrons, which carry the negative charge
Which covalent bonds are polar?
• Bonds between atoms that differ in electronegativity (affinity for electrons)H 2.1 N 3.0 C 2.5 O 3.4
• A bond between atoms that differ by 0.5 - 2.0 is a polar bond. Examples:
O(-)─H(+) N(-)─H(+) C(+)=O(-)
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Polar & ionic bonds
• Electronegativity difference0 - 0.5….………...0.5 - 2………..>2
• Bond typenon-polar cv....polar covalent……ionic
• Sharing of electrons: Equal………..unequal………very unequal
Oppositely charged ions are attracted to each other electrostatically
Ionic bond
Water
• O-H bonds are polar
• Bond angles place the H atoms on one side of the molecule
• Therefore, the water molecule is polar
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Hydrogen bonds among water molecules
Hydrogen bonds
• hydrogen in polar covalent bonds is attracted to nearby electronegative atoms (O or N)
• weak electrostatic bonds – easily broken
• Very important in biology. Examples:– properties of water– protein folding – DNA and RNA folding
Regarding this table, note how strong covalent bonds are compared to other forces holding molecules together.
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Properties of water
• Cohesion
• Surface tension
• Adhesion to hydrophilic substancese.g. cellulose
• Not to hydrophobic substancese.g. waxes
Figure 3.2 Water transport in plants
Surface tension shapes water on a hydrophobic surface
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Figure 3.3 Walking on water
Water physical phases
Ice crystal structure Liquid water Water vapor
Heat
• random movements of atoms and molecules
• add heat: faster movement, higher temperature (heat energy per molecule)
• no heat = “absolute zero” (-273o Celsius, 0o Kelvin)
• units of heat: calorie, kcal = Calorie, calorie=4.184 Joules
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Water stabilizes temperature
• Specific heat: 1 cal/g ºC
• Heat of fusion: ~80 cal/g released by freezing, absorbed by melting
• Heat of vaporization: ~539 cal/g absorbed by evaporation, released by condensation.
• Water expands as it freezes: ice less dense and floats
Black River at Black Rock, AR 6-7 to 8-27-06
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0 7 14 21 28 35 42 49 56 63 70 77 84 91
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One of my temperature recorders, placed in (very) shallow water in the Black River - can you explain the fluctuations?
Floating ice and the fitness of the environment
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A crystal of NaCl dissolving in water
Water is good solvent for polar or ionized substances
Electrolytes
• Compounds held together by ionic bonds that dissolve in polar solvents
• example: sodium chloride (NaCl)becomes Na+ and Cl-
• electrolytes are the most abundant solutes in body fluids- common ions include Na+ Cl- K+ HCO3
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Water is a weak electrolyte
H3O+ or just H+
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Acid-base relations
• In pure water at 20 oC:• [H2O] = 55.4 M • one molecule in 554 million is dissociated• [H+] = 10-7 M• pH = -log [H+] = 7
• pH is the negative logarithm (base 10) of the hydrogen ion concentration
pH of aqueous solutions
“acidic”higher [H+], lower pH
“basic” lower [H+], higher pH
neutral
bufferstabilizes pH
Bio 121-999-F/12 LEC-2